3.6 AS91392 Demonstrate understanding of equilibrium principles in aqueous systems (5 credits)
SOLUBILITY CONSTANT
NCEA PAST EXAM QUESTIONS
i) Write the equation for the equilibrium present in a saturated solution of the following substances
ii) write the expression for Ks for each of the following
iii) Calculate the solubility (or concentration) of the following substances in a saturated solution, in mol L–1
Additional questions:
In an experiment, a saturated solution was made by dissolving 1.44 × 10–3 g of Ag2CrO4 in water, and making it up to a volume of 50.0 mL.
M (Ag2CrO4) = 332 g mol–1. Calculate the solubility of Ag2CrO4(s), and hence give the [Ag+] and [CrO42–] in the solution.Determine the Ks(Ag2CrO4).
Solid sodium chloride is added to 5.00 L of 0.100 mol L–1 silver nitrate solution. Calculate the minimum mass of sodium chloride that would be needed to produce a saturated solution of AgCl. Assume that there is no change in volume when the sodium chloride is added. M(NaCl) = 58.5 g mol–1
a) Describe what is meant by the term ‘solubility’.
b) The chloride ion concentration in sea water can be determined by titrating a sample with aqueous silver nitrate (AgNO3) using potassium chromate (K2CrO4) as the indicator. As the silver nitrate is added, a precipitate of silver chloride, (AgCl) forms. When most of the AgCl has precipitated, the Ag+(aq) concentration becomes high enough for a red precipitate of Ag2CrO4 to form. Show that the solubility of Ag2CrO4 in pure water at 25°C is higher than that of AgCl. Ks(AgCl) = 1.56 × 10–10 Ks(Ag2CrO4) = 1.30 × 10–12
c) If the concentration of chromate ions is 6.30 × 10–3 mol L–1 at the point when the Ag2CrO4 starts to precipitate, calculate the concentration of Ag+ ions in the solution.
QUALITATIVE DESCRIPTIONS AND CALCULATIONS OF SPARINGLY SOLUBLE IONIC SOLIDS
NCEA PAST EXAM QUESTION: Solubility of solids in solutions with a common ion
A sample of seawater has a chloride ion concentration of 0.440 mol L–1. Determine whether a precipitate of lead(II) chloride will form when a 2.00 g sample of lead(II) nitrate is added to 500 mL of the seawater. Ks(PbCl2) = 1.70 × 10–5M(Pb(NO3)2) = 331 g mol–1
Determine whether a precipitate of iron(III) hydroxide, Fe(OH)3, will form when Fe(NO3)3 is dissolved in water. [Fe(NO3)3] = 1.05 × 10–4 mol L–1. Assume the pH of the water is 7. Ks(Fe(OH)3) = 2.00 × 10–39
Discuss how the solubility of Ag2CrO4will change if it is dissolved in 0.1 mol L–1K2CrO4No calculationsare necessary.
Sea water contains many dissolved salts. The chloride ion concentration in a sample of sea water is 0.440 mol L–1. Determine whether a precipitate of lead(II)chloride will form when a 1.00 g sample of lead(II) nitrate is added to 500 mL of the sea water. Your answer must be clearly justified. M(Pb(NO3)2) = 331 g mol–1
Sea-water contains appreciable amounts of ions other than Na+ and Cl–. Evaporating the sea-water to dryness would produce a mixture of salts including NaCl. However, precipitation of NaCl occurs if concentrated hydrochloric acid is added to a saturated NaCl solution. Explain why this precipitation occurs.
Fluoridation of a water supply produces a fluoride concentration of approximately 5 x 10–5 mol L–1.
Will calcium fluoride (CaF2) precipitate in a hard water supply where the concentration of calcium ions is 2 x 10–4 mol L–1? Ks (CaF2)= 3.2 x 10–11
NCEA PAST EXAM QUESTION: Solubility of solids in solutions forming a complex ion
In another experiment, 0.0100 g of Ag2CrO4in beaker A was made up to a volume of 50.0 mL with water. In beaker B, 0.0100 g of Ag2CrO4was made up to a volume of 50.0 mL with 0.100 mol L–1ammonia solution.
Compare and contrast the solubility of Ag2CrO4in beaker A and beaker B. No calculations are necessary.
Solid sodium chloride is added to 5.00 L of 0.100 mol L–1silver nitrate solution. Calculate the minimum mass of sodium chloride that would be needed to produce a saturated solution of AgCl. Assume that there is no change in volume when the sodium chloride is added. M(NaCl) = 58.5 g mol–1
Discuss reasons for the fact that a precipitate of silver chloride dissolves on the addition of excess aqueous ammonia.
NCEA PAST EXAM QUESTION: Solubility of solids in solutions with a change in pH
The solubility of zinc hydroxide, Zn(OH)2, can be altered by changes in pH. Some changes in pH may lead to the formation of complex ions, such as the zincate ion, [Zn(OH)4]2–Use equilibrium principles to explain why the solubility of zinc hydroxide increases when the pH is less than 4 or greater than 10.
No calculations are necessary.
1) Some sulfides have very low solubility products. When hydrogen sulfide gas is bubbled through solutions of these ions, these ions separate from a mixture of ions. i) In a saturated solution of hydrogen sulfide [H3O+]2[S2–] = 1.10 x 10–23Calculate the sulfide ion concentration when the pH of the solution is 4.20.
ii) Calculate the solubility of FeS in this solution, in mol L–1. Ks(FeS) = 4.90 x 10–18
2) A solution contains a mixture of the two metal ions Cu2+ and Zn2+, both of the same concentration. The solution is saturated with hydrogen sulfide and adding hydrochloric acid lowers the pH of the solution. Ks(CuS) = 6.30 x 10–36 Ks(ZnS) = 1.6 x 10–24
Account for the fact that at a pH close to 7 all the metal sulfides will precipitate whereas only the most insoluble sulfides precipitate out at a lower pH.
In your answer, you should use equilibrium principles and both Cu2+ and Zn2+ as examples.(No calculations are required.)
Discuss the effect of decreasing the pH of the water on the solubility of Fe(OH)3.
A saturated solution of zinc hydroxide, Zn(OH)2 contains a small amount of solid Zn(OH)2 at the bottom of the container. The pH of the solution is
increased. Discuss the effect of increasing the pH on the amount of solid present, and also on the nature and concentration of the species present in the
solution. No calculations are necessary.
Discuss how the solubility of Ag2CrO4will change if it is dissolved in 0.1 mol L–1NH3No calculations are necessary.
The Ks of aluminium hydroxide, Al(OH)3, at 25°C, is 3 × 10–34, indicating that it has very low solubility. The solubility may be altered by changes in pH (due to acidic or basic properties) and formation of complex ions such as the aluminate ion, [Al(OH)4]–. Discuss why aluminium hydroxide becomes more soluble in aqueous solutions that have a pH less than 4, or a pH greater than 10. In your answer include:
the equation for the reaction that relates to Ks(Al(OH)3)
equations for the reactions that relate to changes in the solubility of aluminium hydroxide at pH less than 4 or greater than 10
a discussion of the equilibrium principles involved.
An aqueous ammonia solution has a pH of 10 and when phenolphthalein indicator is added it turns pink. Solid ammonium chloride is added to this solution and the solution turns colourless due to a decrease in pH. By considering the equilibrium systems, discuss why the pH of the solution decreased. Include a relevant equation in your answer.
Sea-water contains appreciable amounts of ions other than Na+ and Cl–. As part of the process for extracting table salt from sea-water, sodium hydroxide is added to the sea-water to precipitate the magnesium ions as magnesium hydroxide. The concentration of Mg2+ ions at this stage is 0.555 mol L–1. Calculate the minimum hydroxide ion concentration and hence the pH of the solution needed for precipitation to occur.
NCEA PAST EXAM QUESTION: Relative concentrations of species in solution
When chlorine gas is added to water, the equation for the reaction is:
Cl2(g) + H2O(l) ⇌ HCl(aq) + HOCl(aq)
(a) (i) Write an equation for the reaction of the weak acid, hypochlorous acid, HOCl, with water.
(ii) List all the species present when HOCl reacts with water, in order of decreasing concentration.Justify your order.
1 mol of each of the following substances was placed in separate flasks, and water was added to these flasks to give a total volume of 1 L for each solution, rank these solutions in order of increasing pH. Justify your choice and include equations where appropriate.
CH3NH3Cl
CH3NH2
HCl
The conductivity of the 1 mol L–1solutions formed in (a) can be measured, rank these solutions in order of decreasing conductivity. Compare and contrast the conductivity of each of the 1 mol L–1solutions, with reference to species in solution.
a) i) Write an equation for the reaction of methanoic acid with water.
ii) Methanoic acid, HCOOH, is a weak acid. A dilute aqueous solution of this acid has a pH of 2.78. List all the species in the aqueous solution of methanoic acid in order of decreasing concentration.Give reasons for you answer.
b) Justify the variation in the properties (pH and conductivity) for the four dilute aqueous solutions described in the table below.
(a) Write equations for the reactions occurring when each of the following is added to water.
(i) HCl
(ii) CH3NH2
(iii) NH4Cl
(b) For each of the following 0.100 mol L–1solutions, list all species in order of decreasing concentration.Do not include water
(i) HCl
(ii) CH3NH2
(iii) NH4Cl
(c) Compare and contrast the pH and electrical conductivity of 0.100 mol L–1solutions of HCl, CH3NH2 and NH4Cl. No calculations are necessary.
a) Classify the following 0.100 mol L–1solutions by writing the correct description from the terms below.
strong acid weak acid neutral weak base strong base
NH3 NaCl NH4Cl HF
b) Discuss the relative concentrations of the species present in each of the 0.100 mol L–1solutions of NH3 and HF. You do not need to include water.
Include in your answer:
• any relevant equations
• a ranking of the species present in each solution in order of decreasing concentration
• justification for the ranking of the species. No calculations are necessary.
a) An aqueous solution of ammonium chloride (NH4Cl) has a pH of 4.66.
(i) Write the equation for solid ammonium chloride dissolving in water.
(ii) Write the equation for the ammonium ion reacting with water.
(b) The bar chart below shows the relative concentrations of the species (excluding water) in a solution of 0.1 mol L–1NH4Cl. (The bar chart is not drawn to scale.)
Identify the species A to E. Justify your answer.
The following table lists some properties of aqueous solutions of sodium hydroxide, methylamine and methylammonium chloride.
(a) The solutions above were prepared by adding the compounds to water.
Write equations for the reactions occurring when each of the three compounds are added to water.
NaOH(s)
CH3NH2(g)
CH3NH3Cl(s)
(b) Justify the differences in the pH and conductivity of the three solutions.
When bromine is added to water, it forms hypobromous acid (HOBr), a weak acid. Write an equation to show the equilibrium system that is formed with hypobromous acid and water.
(a) (i) For each of the following 0.1 mol L–1 solutions, write an equation to show the reaction with water.
CH3NH2
NH4Cl
(ii) List all the species in each of the following 0.1 mol L–1 aqueous solutions in order of decreasing concentration. Do not include H2O.
CH3NH2
NH4Cl
(b) Explain why aqueous aminomethane, CH3NH2, is a weak electrolyte.
The boxes below show particle representations of the species (excluding water) in four aqueous solutions
Choose the box that best illustrates each of the solutions (i)–(iii) below. In each case, give a reason for your answer.
i) A dilute solution of a strong acid
ii) A concentrated solution of a weak acid
iii) A buffer solution
Arrange the following 0.1 mol L–1 solutions in order of increasing pH.
NH3
NH4Cl
HCl
NaCl
NaOH
Give reasons for arranging in this order, including equations for any reactions occurring to produce solutions that do not have a pH of 7.
NCEA PAST EXAM QUESTION: Acid Base strength Ka (pKa)
1. Hypochlorous acid has a pKa of 7.53. Another weak acid, hydrofluoric acid, HF, has a pKa of 3.17. A 0.100 mol L–1 solution of each acid was prepared by dissolving it in water. Compare the pHs of these two solutions. No calculations are necessary.
2. Aqueous methylamine, CH3NH2, solution has a pH of 11.8. Show by calculation that the concentration of this solution is 0.0912 mol L–1.
1. 1 mol of each of the following substances was placed in separate flasks, and water was added to these flasks to give a total volume of 1 L for each solution. In the box below, rank these solutions in order of increasing pH. Justify your choice and include equations where appropriate.
CH3NH3Cl
CH3NH2
HCl
2. What is the pH of 20.0 mL of 0.0896 mol L–1 ethanoic acid pKa (CH3COOH) = 4.76
Calculate the concentration of methanoic acid solution with a pH of 2.78. pKa (HCOOH) = 3.74
Calculate the pH of 0.150 mol L–1aqueous ammonia, NH3.pKa (NH4+) = 9.24
1. A solution prepared by dissolving hydrogen fluoride in water has a pH of 2.34. Calculate the concentration of the hydrogen fluoride in the solution. pKa (HF) = 3.17
2. Glycolic acid, HOCH2COOH, is a monoprotic acid used in various skin-care products, and can be represented as HG. Glycolic acid has a pKa value of 3.83.
(a) Write an equation for the reaction of glycolic acid, HG, with water.
(b) Write the Ka expression for glycolic acid, HG.
(c) Calculate the pH of a 0.675 mol L–1 solution of glycolic acid, HG.
An aqueous solution of ammonium chloride (NH4Cl) has a pH of 4.66. Calculate the concentration of NH4Cl solution. pKa(NH4+) = 9.24 Ka = 5.75 × 10 –10
Ethanoic acid, CH3COOH, is a common organic acid. pKa (CH3COOH) = 4.76 Ka = 1.74 × 10–5
(i) Write an equation for the reaction of ethanoic acid with water.
(ii) Write the Ka expression for ethanoic acid.
(iii) Calculate the pH of a 0.0500 mol L–1ethanoic acid solution.
The pH of hydrazoic acid (HN3) is 2.6. Calculate the concentration of the HN3 solution. pKa(HN3) = 4.72
When bromine is added to water, it forms hypobromous acid (HOBr), a weak acid.
(i) Write the Ka expression for hypobromous acid.
(ii) Calculate the pH of a 0.0525 mol L–1 hypobromous acid solution. pKa(HOBr) = 8.62
1. The pH of the solution in the stomach of a patient in hospital is 2.50. As a treatment, the patient is given a small volume of sodium citrate (Na3Cit) solution. Citric acid, H3Cit, is a triprotic acid.
(a) (i) Would the pH of a solution of sodium citrate be less than, equal to or greater than 7? A calculation is not required
(ii) Explain your choice, including an appropriate equation in your answer
2. An aqueous ammonia solution has a pH of 10 and when phenolphthalein indicator is added it turns pink. Solid ammonium chloride is added to this solution and the solution turns colourless due to a decrease in pH. By considering the equilibrium systems, discuss why the pH of the solution decreased. Include a relevant equation in your answer.
NCEA PAST EXAM QUESTION: Titration curves to represent an acid-base system
A titration was carried out by adding hydrobromic acid, HBr, to 20.0 mL of aqueous methylamine, CH3NH2, solution.
The equation for the reaction is: CH3NH2 + HBr → CH3NH3+ + Br– Ka(CH3NH3+) = 2.29 × 10–11 The curve for this titration is given below:
(a) Explain why the pH does not change significantly between the addition of 5 to 15 mL of HBr (around point A on the curve).
Include any relevant equation(s) in your answer.
(b) The aqueous methylamine, CH3NH2, solution has a pH of 11.8 before any HBr is added.
Show by calculation that the concentration of this solution is 0.0912 mol L–1.
(c) (i) Write the formulae of the four chemical species, apart from water and OH–, that are present at the point marked B on the curve.
(ii) Compare and contrast the solution at point B with the initial aqueous methylamine solution.
In your answer you should include:
• a comparison of species present AND their relative concentrations
• a comparison of electrical conductivity linked to the relevant species present in each solution
• equations to support your answer.
20.0 mL of 0.0896 mol L–1 ethanoic acid is titrated with 0.100 mol L–1 sodium hydroxide. pKa (CH3COOH) = 4.76
(a) Calculate the pH of the ethanoic acid before any NaOH is added.
(b) Halfway to the equivalence point of the titration, the pH = pKa of the ethanoic acid.
Discuss the reason for this.
(c) (i) Discuss the change in the concentration of species in solution, as the first 5.00 mL of NaOH is added to the 20.0 mL of ethanoic acid.
Your answer should include chemical equations. No calculations are required.
(ii) Calculate the pH of the titration mixture after 5.00 mL of NaOH has been added.
20.00 mL of 0.125 mol L–1 ethanoic acid is titrated with 0.125 mol L–1 sodium hydroxide solution. The equation for this reaction is:
CH3COOH(aq) + NaOH(aq) --> CH3COONa(aq) + H2O (l)
The titration curve for the reaction is given below and the buffer region is marked on the graph.
a) i) Explain why the solution in the titration flask can act as a buffer in this marked region. Use an equation in your answer.
ii) Put an X on the graph to show at which point the buffering action is the most efficient. Give reasons for your answer.
b) i) Show that the pH at the equivalence point for this titration is 8.78.
pKa(CH3COOH ) = 9.24
ii) Explain why methyl orange is not a suitable indicator for this titration and why phenolphthalein is a suitable indicator for this titration.
iii) Phenolphthalein is an acid-base indicator. It is a weak acid and its formula can be represented as HIn. Phenolphthalein is colourless in acidic solutions and purple in basic solutions.
pKa (HIn) = 9.60
Discuss the effect of adding ethanoic acid and sodium hydroxide in turn to a solution containing phenolphthalein. In your answer, you should refer to:
- equilibrium principles
- the species responsible for the colours seen
- the pH range within which this indicator is effective.
A titration was carried out with methanoic acid and sodium hydroxide. The equation for the reaction is:
HCOOH + NaOH → HCOONa + H2O pKa (HCOOH) = 3.74
The curve for this titration is given below:
25.0 mL of methanoic acid solution is titrated with 0.180 mol L–1sodium hydroxide.
a) i) Show that the concentration of the HCOOH solution is 0.288 mol L–1.
ii)Calculate the initial pH of the 0.288 mol L–1 HCOOH solution.
b) Discuss the pH of the reaction mixture, in terms of the species present, after 20 mL of NaOH has been added. No calculations are necessary.
c) Some indicators and their pKa values are shown in the table below.
Discuss the suitability of each of these indicators for this titration. In your answer you should include:
• an identification of the most suitable indicator(s)
• the consequences of choosing an unsuitable indicator
• an explanation of the significance of the pKa in selecting an indicator.
Below is the titration curve for 10.0 mL of 0.100 mol L–1ethanoic acid being titrated with 0.100 mol L–1sodium hydroxide. Ethanoic acid can be represented by the symbol HEt.
a) With reference to the point marked A on the graph, discuss:
• the species present, and their relative concentrations
• an estimate of the pKa value for ethanoic acid
• the effect of adding small amounts of strong acid or strong base to the solution.
Include relevant equations in your answer. No calculations are necessary.
b) With reference to the point marked B on the graph, discuss the species present, and their effect on the pH at the equivalence point.
Include relevant equations in your answer. No calculations are necessary.
20.00 mL of 0.160 mol L–1 ammonia is titrated with 0.230 mol L–1 hydrochloric acid. The equation for the reaction is
a) Explain, in terms of species present, why the pH at B (half way to the equivalence volume) is 9.24.
b) Calculate the pH at point A.
c) Discuss the pH of the reaction mixture at point C, in terms of the species present.
25.0 mL of 0.0500 mol L–1benzoic acid solution (C6H5COOH) is titrated with 0.0500 mol L–1sodium hydroxide solution. The equation for the reaction is:
C6H5COOH(aq) + NaOH(aq) → C6H5COONa(aq) + H2O(ℓ)
The titration curve for the reaction is:
a) Write the formulae of the four chemical species, apart from water and H3O+, that are present at the equivalence point.
b) Explain why the solution in the titration flask has buffering properties after 9.80 mL of the NaOH solution has been added, but not when 25.0 mL has been added.
c) Some indicators are shown in the table below.
Discuss the suitability of these indicators for this titration. Your discussion should include:
• identification of the most suitable indicator(s)
• consideration of how indicators are chosen for a titration
• the consequences of choosing an unsuitable indicator.
The following titration curve shows the addition of aqueous 0.100 mol L–1sodium hydroxide to a solution of hydrazoic acid, HN3.
pKa(HN3) = 4.72
a) i) Draw a cross (X) on the titration curve to indicate the pH at the equivalence point of the titration.
ii) Complete the titration curve to show how the pH changes as more aqueous sodium hydroxide is added.
b) The initial pH of the hydrazoic acid (HN3) is 2.6. Calculate the concentration of the HN3solution used in the titration.
7) A 0.160 mol L–1 solution of sodium hydroxide is titrated against 50 mL of aqueous propanoic acid, HPr. 40 mL of the sodium hydroxide solution was required to exactly react with the propanoic acid. The reaction occurring can be represented as: HPr(aq) + NaOH(aq) → NaPr(aq) + H2O Ka(HPr) = 1.35 × 10–5
a) i) Show that the concentration of the aqueous propanoic acid is 0.128 mol L–1.
ii) Calculate the pH of the aqueous propanoic acid.
b) Calculate the pH at the equivalence point.
c) 35 mL of the sodium hydroxide solution is added to a second 50 mL sample of the same acid to form a buffer solution.
i) What is the function of a buffer?
ii) Discuss the ability of the solution formed to act as a buffer. Your answer should include relevant equations.
d) The equivalence point of the titration could also be found using an acid-base indicator. Which of the following indicators would be suitable to use? Explain your choice of indicator.
The graph below shows the change in pH when 40.0 mL of 0.0500 mol L–1 aqueous NH3 is titrated with 0.200 mol L–1 aqueous HCl
The equation for the reaction occurring during the titration is: NH3(aq) + HCl(aq) → NH4Cl(aq)
a) Use the curve to determine pKa(NH4+) and hence calculate Ka(NH4+).
b) Explain why the pH at the equivalence point for this titration is less than 7. (Include an equation to support your answer.)
A NH4+ /NH3 buffer solution is prepared with a pH of 9.6.
c) Use the graph to describe how this buffer solution could be made from 0.0500 mol L–1 NH3 solution and 0.200 mol L–1 HCl solutions.
A second titration is carried out – this time 40.0 mL of 0.0500 mol L–1 NH4Cl solution is titrated against 0.200 mol L–1 NaOH solution.
d) Write an equation for the titration reaction.
e) i) Show that [NH3] at the equivalence point is 0.0400 mol L–1.
ii) Using Ka(NH4+) determined in part (a) on the previous page, determine the pH at the equivalence point of the second titration.
The active ingredient in many sunscreens is para-aminobenzoic acid. It is a weak monoprotic acid and can be represented as HPab, while its conjugate base is Pab–.
a) Write an equation for the reactions occurring at equilibrium when HPab is dissolved in water.
b) Write the expression for Ka(HPab).
A solution of HPab in water was prepared at 25oC and its pH was found to be 3.22.
c) Calculate the concentration of H3O+ in the solution.
d) The concentration of the HPab solution was determined by titration. A 20.0 mL sample of the HPab solution required 12.0 mL of 0.0500 mol L–1 NaOH to reach the equivalence point. The
equation for the reaction occurring is HPab + NaOH → NaPab + H2O
i) Calculate the concentration of the HPab solution
ii) Using the results from parts (c) and (d)(i), show that pKa(HPab) = 4.92.
e) Would the pH at the equivalence point of the titration of HPab with NaOH be more than 7, less than 7 or equal to 7? Give reasons and include any relevant equations that support your answer.
f) Using the information above, sketch a curve showing the change in pH against the volume of sodium hydroxide added to the 20.0 mL HPab solution in the flask.
An aqueous solution containing a mixture of HF and sodium fluoride, NaF, can act as a buffer solution. Calculate the mass of NaF that must be added to 150 mL of 0.0500 mol L–1 HF to give a buffer solution with a pH of 4.02. Assume there is no change in volume. M(NaF) = 42.0 g mol–1 pKa(HF) = 3.17
(i) The following two solutions are mixed to form a buffer solution: 20.0 mL of 1 mol L–1
CH3NH3Cl and 30.0 mL of 1 mol L–1 CH3NH2
Calculate the pH of the resultant buffer solution. pKa (CH3NH3+) = 10.64
(ii) Explain the effect on the solution formed in (i) when a small amount of acid is added
a) A mixture of aqueous solutions of NH3and ammonium chloride, NH4Cl, can act as a buffer solution. Calculate the mass of NH4Cl required, when added to 250 mL of a 0.150 mol L–1NH3solution, to give a buffer solution with a pH of 8.60. Assume there is no change in volume.
M (NH4Cl) = 53.5 g mol–1 pKa (NH4+) = 9.24
b) Discuss the ability of the NH3 / NH4Cl solution to act as a buffer at a pH of 8.60. In you answer you should: • describe the function of a buffer solution • evaluate its effectiveness when small amounts of acid or base are added • include any relevant equations.
A buffer solution is made by adding solid sodium methanoate, HCOONa, to an aqueous solution of methanoic acid, HCOOH. pKa(HCOOH) = 3.74
a) Describe the function of a buffer solution.
b) Explain why the solution made with methanoic acid, HCOOH, and sodium methanoate, HCOONa, has the ability to act as a buffer.
3.6 AS91392 Demonstrate understanding of equilibrium principles in aqueous systems (5 credits)
SOLUBILITY CONSTANT
NCEA PAST EXAM QUESTIONS
i) Write the equation for the equilibrium present in a saturated solution of the following substances
ii) write the expression for Ks for each of the following
iii) Calculate the solubility (or concentration) of the following substances in a saturated solution, in mol L–1
Additional questions:
In an experiment, a saturated solution was made by dissolving 1.44 × 10–3 g of Ag2CrO4 in water, and making it up to a volume of 50.0 mL.
M (Ag2CrO4) = 332 g mol–1. Calculate the solubility of Ag2CrO4(s), and hence give the [Ag+] and [CrO42–] in the solution.Determine the Ks(Ag2CrO4).
Solid sodium chloride is added to 5.00 L of 0.100 mol L–1 silver nitrate solution. Calculate the minimum mass of sodium chloride that would be needed to produce a saturated solution of AgCl. Assume that there is no change in volume when the sodium chloride is added. M(NaCl) = 58.5 g mol–1
a) Describe what is meant by the term ‘solubility’.
b) The chloride ion concentration in sea water can be determined by titrating a sample with aqueous silver nitrate (AgNO3) using potassium chromate (K2CrO4) as the indicator. As the silver nitrate is added, a precipitate of silver chloride, (AgCl) forms. When most of the AgCl has precipitated, the Ag+(aq) concentration becomes high enough for a red precipitate of Ag2CrO4 to form. Show that the solubility of Ag2CrO4 in pure water at 25°C is higher than that of AgCl. Ks(AgCl) = 1.56 × 10–10 Ks(Ag2CrO4) = 1.30 × 10–12
c) If the concentration of chromate ions is 6.30 × 10–3 mol L–1 at the point when the Ag2CrO4 starts to precipitate, calculate the concentration of Ag+ ions in the solution.
QUALITATIVE DESCRIPTIONS AND CALCULATIONS OF SPARINGLY SOLUBLE IONIC SOLIDS
NCEA PAST EXAM QUESTION: Solubility of solids in solutions with a common ion
A sample of seawater has a chloride ion concentration of 0.440 mol L–1. Determine whether a precipitate of lead(II) chloride will form when a 2.00 g sample of lead(II) nitrate is added to 500 mL of the seawater. Ks(PbCl2) = 1.70 × 10–5M(Pb(NO3)2) = 331 g mol–1
Determine whether a precipitate of iron(III) hydroxide, Fe(OH)3, will form when Fe(NO3)3 is dissolved in water. [Fe(NO3)3] = 1.05 × 10–4 mol L–1. Assume the pH of the water is 7. Ks(Fe(OH)3) = 2.00 × 10–39
Discuss how the solubility of Ag2CrO4will change if it is dissolved in 0.1 mol L–1K2CrO4No calculationsare necessary.
Sea water contains many dissolved salts. The chloride ion concentration in a sample of sea water is 0.440 mol L–1. Determine whether a precipitate of lead(II)chloride will form when a 1.00 g sample of lead(II) nitrate is added to 500 mL of the sea water. Your answer must be clearly justified. M(Pb(NO3)2) = 331 g mol–1
Sea-water contains appreciable amounts of ions other than Na+ and Cl–. Evaporating the sea-water to dryness would produce a mixture of salts including NaCl. However, precipitation of NaCl occurs if concentrated hydrochloric acid is added to a saturated NaCl solution. Explain why this precipitation occurs.
Fluoridation of a water supply produces a fluoride concentration of approximately 5 x 10–5 mol L–1.
Will calcium fluoride (CaF2) precipitate in a hard water supply where the concentration of calcium ions is 2 x 10–4 mol L–1? Ks (CaF2)= 3.2 x 10–11
NCEA PAST EXAM QUESTION: Solubility of solids in solutions forming a complex ion
In another experiment, 0.0100 g of Ag2CrO4in beaker A was made up to a volume of 50.0 mL with water. In beaker B, 0.0100 g of Ag2CrO4was made up to a volume of 50.0 mL with 0.100 mol L–1ammonia solution.
Compare and contrast the solubility of Ag2CrO4in beaker A and beaker B. No calculations are necessary.
Discuss reasons for the fact that a precipitate of silver chloride dissolves on the addition of excess aqueous ammonia.
NCEA PAST EXAM QUESTION: Solubility of solids in solutions with a change in pH
The solubility of zinc hydroxide, Zn(OH)2, can be altered by changes in pH. Some changes in pH may lead to the formation of complex ions, such as the zincate ion, [Zn(OH)4]2– Use equilibrium principles to explain why the solubility of zinc hydroxide increases when the pH is less than 4 or greater than 10.
No calculations are necessary.
1) Some sulfides have very low solubility products. When hydrogen sulfide gas is bubbled through solutions of these ions, these ions separate from a mixture of ions. i) In a saturated solution of hydrogen sulfide [H3O+]2[S2–] = 1.10 x 10–23Calculate the sulfide ion concentration when the pH of the solution is 4.20.
ii) Calculate the solubility of FeS in this solution, in mol L–1. Ks(FeS) = 4.90 x 10–18
2) A solution contains a mixture of the two metal ions Cu2+ and Zn2+, both of the same concentration. The solution is saturated with hydrogen sulfide and adding hydrochloric acid lowers the pH of the solution. Ks(CuS) = 6.30 x 10–36 Ks(ZnS) = 1.6 x 10–24
Account for the fact that at a pH close to 7 all the metal sulfides will precipitate whereas only the most insoluble sulfides precipitate out at a lower pH.
In your answer, you should use equilibrium principles and both Cu2+ and Zn2+ as examples.(No calculations are required.)
Discuss the effect of decreasing the pH of the water on the solubility of Fe(OH)3.
A saturated solution of zinc hydroxide, Zn(OH)2 contains a small amount of solid Zn(OH)2 at the bottom of the container. The pH of the solution is
increased. Discuss the effect of increasing the pH on the amount of solid present, and also on the nature and concentration of the species present in the
solution. No calculations are necessary.
Discuss how the solubility of Ag2CrO4will change if it is dissolved in 0.1 mol L–1NH3No calculations are necessary.
The Ks of aluminium hydroxide, Al(OH)3, at 25°C, is 3 × 10–34, indicating that it has very low solubility. The solubility may be altered by changes in pH (due to acidic or basic properties) and formation of complex ions such as the aluminate ion, [Al(OH)4]–. Discuss why aluminium hydroxide becomes more soluble in aqueous solutions that have a pH less than 4, or a pH greater than 10. In your answer include:
the equation for the reaction that relates to Ks(Al(OH)3)
equations for the reactions that relate to changes in the solubility of aluminium hydroxide at pH less than 4 or greater than 10
a discussion of the equilibrium principles involved.
An aqueous ammonia solution has a pH of 10 and when phenolphthalein indicator is added it turns pink. Solid ammonium chloride is added to this solution and the solution turns colourless due to a decrease in pH. By considering the equilibrium systems, discuss why the pH of the solution decreased. Include a relevant equation in your answer.
Sea-water contains appreciable amounts of ions other than Na+ and Cl–. As part of the process for extracting table salt from sea-water, sodium hydroxide is added to the sea-water to precipitate the magnesium ions as magnesium hydroxide. The concentration of Mg2+ ions at this stage is 0.555 mol L–1. Calculate the minimum hydroxide ion concentration and hence the pH of the solution needed for precipitation to occur.
Ks (Mg(OH)2) = 7.10 × 10–12
Questions & Answers from the old Unit Standard task "Will a precipitate form"
RELATIVE CONCENTRATIONS OF SPECIES IN SOLUTION
NCEA PAST EXAM QUESTION: Relative concentrations of species in solution
When chlorine gas is added to water, the equation for the reaction is:
Cl2(g) + H2O(l) ⇌ HCl(aq) + HOCl(aq)(a) (i) Write an equation for the reaction of the weak acid, hypochlorous acid, HOCl, with water.
(ii) List all the species present when HOCl reacts with water, in order of decreasing concentration.Justify your order.
1 mol of each of the following substances was placed in separate flasks, and water was added to these flasks to give a total volume of 1 L for each solution, rank these solutions in order of increasing pH. Justify your choice and include equations where appropriate.
CH3NH3Cl
CH3NH2
HCl
The conductivity of the 1 mol L–1solutions formed in (a) can be measured, rank these solutions in order of decreasing conductivity. Compare and contrast the conductivity of each of the 1 mol L–1solutions, with reference to species in solution.
a) i) Write an equation for the reaction of methanoic acid with water.
ii) Methanoic acid, HCOOH, is a weak acid. A dilute aqueous solution of this acid has a pH of 2.78. List all the species in the aqueous solution of methanoic acid in order of decreasing concentration.Give reasons for you answer.
b) Justify the variation in the properties (pH and conductivity) for the four dilute aqueous solutions described in the table below.
(a) Write equations for the reactions occurring when each of the following is added to water.
(i) HCl
(ii) CH3NH2
(iii) NH4Cl
(b) For each of the following 0.100 mol L–1solutions, list all species in order of decreasing concentration.Do not include water
(i) HCl
(ii) CH3NH2
(iii) NH4Cl
(c) Compare and contrast the pH and electrical conductivity of 0.100 mol L–1solutions of HCl, CH3NH2 and NH4Cl. No calculations are necessary.
a) Classify the following 0.100 mol L–1solutions by writing the correct description from the terms below.
strong acid weak acid neutral weak base strong base
NH3 NaCl NH4Cl HF
b) Discuss the relative concentrations of the species present in each of the 0.100 mol L–1solutions of NH3 and HF. You do not need to include water.
Include in your answer:
• any relevant equations
• a ranking of the species present in each solution in order of decreasing concentration
• justification for the ranking of the species. No calculations are necessary.
a) An aqueous solution of ammonium chloride (NH4Cl) has a pH of 4.66.
(i) Write the equation for solid ammonium chloride dissolving in water.
(ii) Write the equation for the ammonium ion reacting with water.
(b) The bar chart below shows the relative concentrations of the species (excluding water) in a solution of 0.1 mol L–1NH4Cl. (The bar chart is not drawn to scale.)
The following table lists some properties of aqueous solutions of sodium hydroxide, methylamine and methylammonium chloride.
(a) The solutions above were prepared by adding the compounds to water.
Write equations for the reactions occurring when each of the three compounds are added to water.
NaOH(s)
CH3NH2(g)
CH3NH3Cl(s)
(b) Justify the differences in the pH and conductivity of the three solutions.
When bromine is added to water, it forms hypobromous acid (HOBr), a weak acid. Write an equation to show the equilibrium system that is formed with hypobromous acid and water.
(a) (i) For each of the following 0.1 mol L–1 solutions, write an equation to show the reaction with water.
CH3NH2
NH4Cl
(ii) List all the species in each of the following 0.1 mol L–1 aqueous solutions in order of decreasing concentration. Do not include H2O.
CH3NH2
NH4Cl
(b) Explain why aqueous aminomethane, CH3NH2, is a weak electrolyte.
The boxes below show particle representations of the species (excluding water) in four aqueous solutions
Choose the box that best illustrates each of the solutions (i)–(iii) below. In each case, give a reason for your answer.
i) A dilute solution of a strong acid
ii) A concentrated solution of a weak acid
iii) A buffer solution
Arrange the following 0.1 mol L–1 solutions in order of increasing pH.
NH3
NH4Cl
HCl
NaCl
NaOH
Give reasons for arranging in this order, including equations for any reactions occurring to produce solutions that do not have a pH of 7.
ACID BASE STRENGTH Ka (pKa)
NCEA PAST EXAM QUESTION: Acid Base strength Ka (pKa)
1. Hypochlorous acid has a pKa of 7.53. Another weak acid, hydrofluoric acid, HF, has a pKa of 3.17. A 0.100 mol L–1 solution of each acid was prepared by dissolving it in water. Compare the pHs of these two solutions. No calculations are necessary.
2. Aqueous methylamine, CH3NH2, solution has a pH of 11.8. Show by calculation that the concentration of this solution is 0.0912 mol L–1.
1. 1 mol of each of the following substances was placed in separate flasks, and water was added to these flasks to give a total volume of 1 L for each solution. In the box below, rank these solutions in order of increasing pH. Justify your choice and include equations where appropriate.
CH3NH3Cl
CH3NH2
HCl
2. What is the pH of 20.0 mL of 0.0896 mol L–1 ethanoic acid pKa (CH3COOH) = 4.76
Calculate the concentration of methanoic acid solution with a pH of 2.78. pKa (HCOOH) = 3.74
Calculate the pH of 0.150 mol L–1aqueous ammonia, NH3.pKa (NH4+) = 9.24
1. A solution prepared by dissolving hydrogen fluoride in water has a pH of 2.34. Calculate the concentration of the hydrogen fluoride in the solution. pKa (HF) = 3.17
2. Glycolic acid, HOCH2COOH, is a monoprotic acid used in various skin-care products, and can be represented as HG. Glycolic acid has a pKa value of 3.83.
(a) Write an equation for the reaction of glycolic acid, HG, with water.
(b) Write the Ka expression for glycolic acid, HG.
(c) Calculate the pH of a 0.675 mol L–1 solution of glycolic acid, HG.
An aqueous solution of ammonium chloride (NH4Cl) has a pH of 4.66. Calculate the concentration of NH4Cl solution. pKa(NH4+) = 9.24 Ka = 5.75 × 10 –10
Ethanoic acid, CH3COOH, is a common organic acid. pKa (CH3COOH) = 4.76 Ka = 1.74 × 10–5
(i) Write an equation for the reaction of ethanoic acid with water.
(ii) Write the Ka expression for ethanoic acid.
(iii) Calculate the pH of a 0.0500 mol L–1ethanoic acid solution.
The pH of hydrazoic acid (HN3) is 2.6. Calculate the concentration of the HN3 solution. pKa(HN3) = 4.72
When bromine is added to water, it forms hypobromous acid (HOBr), a weak acid.
(i) Write the Ka expression for hypobromous acid.
(ii) Calculate the pH of a 0.0525 mol L–1 hypobromous acid solution. pKa(HOBr) = 8.62
1. The pH of the solution in the stomach of a patient in hospital is 2.50. As a treatment, the patient is given a small volume of sodium citrate (Na3Cit) solution. Citric acid, H3Cit, is a triprotic acid.
(a) (i) Would the pH of a solution of sodium citrate be less than, equal to or greater than 7? A calculation is not required
(ii) Explain your choice, including an appropriate equation in your answer
2. An aqueous ammonia solution has a pH of 10 and when phenolphthalein indicator is added it turns pink. Solid ammonium chloride is added to this solution and the solution turns colourless due to a decrease in pH. By considering the equilibrium systems, discuss why the pH of the solution decreased. Include a relevant equation in your answer.
Do Task A from this old unit standard task "Acids and Bases"
TITRATION CURVES
NCEA PAST EXAM QUESTION: Titration curves to represent an acid-base system
A titration was carried out by adding hydrobromic acid, HBr, to 20.0 mL of aqueous methylamine, CH3NH2, solution.
The equation for the reaction is: CH3NH2 + HBr → CH3NH3+ + Br– Ka(CH3NH3+) = 2.29 × 10–11 The curve for this titration is given below:
(a) Explain why the pH does not change significantly between the addition of 5 to 15 mL of HBr (around point A on the curve).
Include any relevant equation(s) in your answer.
(b) The aqueous methylamine, CH3NH2, solution has a pH of 11.8 before any HBr is added.
Show by calculation that the concentration of this solution is 0.0912 mol L–1.
(c) (i) Write the formulae of the four chemical species, apart from water and OH–, that are present at the point marked B on the curve.
(ii) Compare and contrast the solution at point B with the initial aqueous methylamine solution.
In your answer you should include:
• a comparison of species present AND their relative concentrations
• a comparison of electrical conductivity linked to the relevant species present in each solution
• equations to support your answer.
20.0 mL of 0.0896 mol L–1 ethanoic acid is titrated with 0.100 mol L–1 sodium hydroxide. pKa (CH3COOH) = 4.76
(a) Calculate the pH of the ethanoic acid before any NaOH is added.
(b) Halfway to the equivalence point of the titration, the pH = pKa of the ethanoic acid.
Discuss the reason for this.
(c) (i) Discuss the change in the concentration of species in solution, as the first 5.00 mL of NaOH is added to the 20.0 mL of ethanoic acid.
Your answer should include chemical equations. No calculations are required.
(ii) Calculate the pH of the titration mixture after 5.00 mL of NaOH has been added.
20.00 mL of 0.125 mol L–1 ethanoic acid is titrated with 0.125 mol L–1 sodium hydroxide solution. The equation for this reaction is:
CH3COOH(aq) + NaOH(aq) --> CH3COONa(aq) + H2O (l)The titration curve for the reaction is given below and the buffer region is marked on the graph.
a) i) Explain why the solution in the titration flask can act as a buffer in this marked region. Use an equation in your answer.
ii) Put an X on the graph to show at which point the buffering action is the most efficient. Give reasons for your answer.
b) i) Show that the pH at the equivalence point for this titration is 8.78.
pKa(CH3COOH ) = 9.24ii) Explain why methyl orange is not a suitable indicator for this titration and why phenolphthalein is a suitable indicator for this titration.
iii) Phenolphthalein is an acid-base indicator. It is a weak acid and its formula can be represented as HIn. Phenolphthalein is colourless in acidic solutions and purple in basic solutions.
pKa (HIn) = 9.60Discuss the effect of adding ethanoic acid and sodium hydroxide in turn to a solution containing phenolphthalein. In your answer, you should refer to:
- equilibrium principles
- the species responsible for the colours seen
- the pH range within which this indicator is effective.
A titration was carried out with methanoic acid and sodium hydroxide. The equation for the reaction is:
HCOOH + NaOH → HCOONa + H2O pKa (HCOOH) = 3.74
The curve for this titration is given below:
25.0 mL of methanoic acid solution is titrated with 0.180 mol L–1sodium hydroxide.
a) i) Show that the concentration of the HCOOH solution is 0.288 mol L–1.
ii) Calculate the initial pH of the 0.288 mol L–1 HCOOH solution.
b) Discuss the pH of the reaction mixture, in terms of the species present, after 20 mL of NaOH has been added. No calculations are necessary.
c) Some indicators and their pKa values are shown in the table below.
Discuss the suitability of each of these indicators for this titration. In your answer you should include:
• an identification of the most suitable indicator(s)
• the consequences of choosing an unsuitable indicator
• an explanation of the significance of the pKa in selecting an indicator.
Below is the titration curve for 10.0 mL of 0.100 mol L–1ethanoic acid being titrated with 0.100 mol L–1sodium hydroxide. Ethanoic acid can be represented by the symbol HEt.
a) With reference to the point marked A on the graph, discuss:
• the species present, and their relative concentrations
• an estimate of the pKa value for ethanoic acid
• the effect of adding small amounts of strong acid or strong base to the solution.
Include relevant equations in your answer. No calculations are necessary.
b) With reference to the point marked B on the graph, discuss the species present, and their effect on the pH at the equivalence point.
Include relevant equations in your answer. No calculations are necessary.
20.00 mL of 0.160 mol L–1 ammonia is titrated with 0.230 mol L–1 hydrochloric acid. The equation for the reaction is
NH3+ HCl → NH4++ Cl–pKa(NH4+) = 9.24, Ka = 5.75 × 10–10The curve for this titration is given below.
a) Explain, in terms of species present, why the pH at B (half way to the equivalence volume) is 9.24.
b) Calculate the pH at point A.
c) Discuss the pH of the reaction mixture at point C, in terms of the species present.
25.0 mL of 0.0500 mol L–1benzoic acid solution (C6H5COOH) is titrated with 0.0500 mol L–1sodium hydroxide solution. The equation for the reaction is:
C6H5COOH(aq) + NaOH(aq) → C6H5COONa(aq) + H2O(ℓ)The titration curve for the reaction is:
a) Write the formulae of the four chemical species, apart from water and H3O+, that are present at the equivalence point.
b) Explain why the solution in the titration flask has buffering properties after 9.80 mL of the NaOH solution has been added, but not when 25.0 mL has been added.
c) Some indicators are shown in the table below.
Discuss the suitability of these indicators for this titration. Your discussion should include:
• identification of the most suitable indicator(s)
• consideration of how indicators are chosen for a titration
• the consequences of choosing an unsuitable indicator.
The following titration curve shows the addition of aqueous 0.100 mol L–1sodium hydroxide to a solution of hydrazoic acid, HN3.
pKa(HN3) = 4.72a) i) Draw a cross (X) on the titration curve to indicate the pH at the equivalence point of the titration.
ii) Complete the titration curve to show how the pH changes as more aqueous sodium hydroxide is added.
b) The initial pH of the hydrazoic acid (HN3) is 2.6. Calculate the concentration of the HN3solution used in the titration.
7) A 0.160 mol L–1 solution of sodium hydroxide is titrated against 50 mL of aqueous propanoic acid, HPr. 40 mL of the sodium hydroxide solution was required to exactly react with the propanoic acid. The reaction occurring can be represented as: HPr(aq) + NaOH(aq) → NaPr(aq) + H2O Ka(HPr) = 1.35 × 10–5
a) i) Show that the concentration of the aqueous propanoic acid is 0.128 mol L–1.
ii) Calculate the pH of the aqueous propanoic acid.
b) Calculate the pH at the equivalence point.
c) 35 mL of the sodium hydroxide solution is added to a second 50 mL sample of the same acid to form a buffer solution.
i) What is the function of a buffer?
ii) Discuss the ability of the solution formed to act as a buffer. Your answer should include relevant equations.
d) The equivalence point of the titration could also be found using an acid-base indicator. Which of the following indicators would be suitable to use? Explain your choice of indicator.
The graph below shows the change in pH when 40.0 mL of 0.0500 mol L–1 aqueous NH3 is titrated with 0.200 mol L–1 aqueous HCl
The equation for the reaction occurring during the titration is: NH3(aq) + HCl(aq) → NH4Cl(aq)b) Explain why the pH at the equivalence point for this titration is less than 7. (Include an equation to support your answer.)
A NH4+ /NH3 buffer solution is prepared with a pH of 9.6.
c) Use the graph to describe how this buffer solution could be made from 0.0500 mol L–1 NH3 solution and 0.200 mol L–1 HCl solutions.
A second titration is carried out – this time 40.0 mL of 0.0500 mol L–1 NH4Cl solution is titrated against 0.200 mol L–1 NaOH solution.
d) Write an equation for the titration reaction.
e) i) Show that [NH3] at the equivalence point is 0.0400 mol L–1.
ii) Using Ka(NH4+) determined in part (a) on the previous page, determine the pH at the equivalence point of the second titration.
The active ingredient in many sunscreens is para-aminobenzoic acid. It is a weak monoprotic acid and can be represented as HPab, while its conjugate base is Pab–.
a) Write an equation for the reactions occurring at equilibrium when HPab is dissolved in water.
b) Write the expression for Ka(HPab).
A solution of HPab in water was prepared at 25oC and its pH was found to be 3.22.
c) Calculate the concentration of H3O+ in the solution.
d) The concentration of the HPab solution was determined by titration. A 20.0 mL sample of the HPab solution required 12.0 mL of 0.0500 mol L–1 NaOH to reach the equivalence point. The
equation for the reaction occurring is HPab + NaOH → NaPab + H2O
i) Calculate the concentration of the HPab solution
ii) Using the results from parts (c) and (d)(i), show that pKa(HPab) = 4.92.
e) Would the pH at the equivalence point of the titration of HPab with NaOH be more than 7, less than 7 or equal to 7? Give reasons and include any relevant equations that support your answer.
f) Using the information above, sketch a curve showing the change in pH against the volume of sodium hydroxide added to the 20.0 mL HPab solution in the flask.
Try this old unit standard task "Acid-Base titration curves"
BUFFER SOLUTIONS
NCEA QUESTIONS ON BUFFER SOLUTIONS
An aqueous solution containing a mixture of HF and sodium fluoride, NaF, can act as a buffer solution. Calculate the mass of NaF that must be added to 150 mL of 0.0500 mol L–1 HF to give a buffer solution with a pH of 4.02. Assume there is no change in volume. M(NaF) = 42.0 g mol–1 pKa(HF) = 3.17
(i) The following two solutions are mixed to form a buffer solution: 20.0 mL of 1 mol L–1
CH3NH3Cl and 30.0 mL of 1 mol L–1 CH3NH2
Calculate the pH of the resultant buffer solution. pKa (CH3NH3+) = 10.64
(ii) Explain the effect on the solution formed in (i) when a small amount of acid is added
a) A mixture of aqueous solutions of NH3and ammonium chloride, NH4Cl, can act as a buffer solution. Calculate the mass of NH4Cl required, when added to 250 mL of a 0.150 mol L–1NH3solution, to give a buffer solution with a pH of 8.60. Assume there is no change in volume.
M (NH4Cl) = 53.5 g mol–1 pKa (NH4+) = 9.24
b) Discuss the ability of the NH3 / NH4Cl solution to act as a buffer at a pH of 8.60. In you answer you should:• describe the function of a buffer solution
• evaluate its effectiveness when small amounts of acid or base are added
• include any relevant equations.
A buffer solution is made by adding solid sodium methanoate, HCOONa, to an aqueous solution of methanoic acid, HCOOH. pKa(HCOOH) = 3.74
a) Describe the function of a buffer solution.
b) Explain why the solution made with methanoic acid, HCOOH, and sodium methanoate, HCOONa, has the ability to act as a buffer.
Your answer should include relevant equations.
Ks (Mg(OH)2) = 7.10 × 10–12